0
To calculate the number of photons, you need the formula E=hf where h is Planck's constant with a value of 6.63*10^-34Js and f should be given.
If f isn't given, then use the formula C = f * wavelength. Rearrange this formula by using the wavelength given and the C, speed of light, which is 3*10^8.
You should get C/wavelength = f, which will then be placed into E=hf => answer.
What you also need is the Intensity. This way you obtain the photon flux as: I/E (i.e. the number of photons per unit area and unit time).
What else can I help you with?
How do you calculate how much energy in kJ do 3.0 moles of photons all with a wavelength of 655 nm contain?
The energy is 18,263.10e4 joules.
How do you calculate energy per mole of photons if you know joules per photon?
To calculate the energy per mole of photons from the energy per photon, you need to multiply the energy per photon by Avogadro's number (6.022 x 10^23) to account for the number of photons in a mole. The formula is: Energy per mole of photons = Energy per photon x Avogadro's number.
How much energy in kJ do 3.0 moles of photons all with a wavelength of 655 nm contain?
To calculate the energy of photons, you can use the equation E = hc/λ, where h is Planck's constant (6.626 x 10^-34 J·s), c is the speed of light (3.00 x 10^8 m/s), and λ is the wavelength. First, convert the wavelength to meters (655 nm = 655 x 10^-9 m). Plug the values into the equation to find the energy per photon, and then multiply by Avogadro's number to get the total energy for 3.0 moles of photons.
How much energy (in kJ) do 3.0 moles of photons all with a wavelength of 675 nm contain?
To calculate the energy of photons, we can use the formula (E = \frac{hc}{\lambda}), where (h) is Planck's constant ((6.626 \times 10^{-34} , \text{J s})), (c) is the speed of light ((3.00 \times 10^8 , \text{m/s})), and (\lambda) is the wavelength in meters (675 nm = (675 \times 10^{-9} , \text{m})). First, calculate the energy of one photon, then multiply by the number of moles (using Avogadro's number, (6.022 \times 10^{23} , \text{photons/mole})). Calculating this gives: Energy of one photon: [ E = \frac{(6.626 \times 10^{-34} , \text{J s})(3.00 \times 10^8 , \text{m/s})}{675 \times 10^{-9} , \text{m}} \approx 2.94 \times 10^{-19} , \text{J} ] Total energy for 3.0 moles of photons: [ \text{Total energy} = 3.0 , \text{moles} \times (6.022 \times 10^{23} , \text{photons/mole}) \times (2.94 \times 10^{-19} , \text{J}) \approx 5.34 \times 10^{5} , \text{J} ] Convert to kJ: [ 5.34 \times 10^{5} , \text{J} \div 1000 \approx 534 , \text{kJ} ] Thus, 3.0 moles of photons at 675 nm contain approximately 534 kJ of energy.
How many moles of photons with a wavelength of 660 nm are needed to provide 6 kJ of energy?
The energy carried by a photon is given byE = hfWhere h is Planck's constant (6.626x10^-34 Joule-seconds) and f is the frequency of the photon in Hertz (Hz).We are given the wavelength of the photon in the question in nanometers (nm). First, we need to convert this to (SI) units, because our equations only work with SI units. Then, we will calculate the frequency of the photon from its wavelength. Once we know the frequency of the photon we're interested in, we simply use the equation above to find the energy carried by one of them. Then we divide 6 kJ by that amount of energy, and the quotient will be the number of photons needed to carry 6 kJ. Finally, when we know the number of photons we need, we divide by the number of photons in a mole to get the number of moles.The SI unit of length is the meter (m). 1nanometer (nm) is 10^-9 meters.660 nm = 660 *10^-9 m = 6.6*10^-7 m.Now we will calculate this photon's frequency from its wavelength. These are related by the equationc = fLwhere c is the speed of light (3*10^8 m/s), f is the frequency of the photon and L is the wavelength of the photon.c = fL(3*10^8 m/s) = f * (6.6*10^-7 m)solving for f, we havef = (3*10^8 m/s) / (6.6*10^-7 m) = 4.54*10^15 s^-1Note that the unit of seconds (s) raised to the -1power is defined as 1 Hertz (Hz).f = 4.54*10^15 HzNow we will use the top equation to solve for the energy carried by one photon having this frequency.E = hfE = (6.626*10^-34 Js) * (4.54*10^15 Hz)E = 1.369*10^-17 JThis is how much energy is carried by one photon of wavelength 660 nm (which will also have a frequency of 4.54*10^15 Hz).How many of these do we need to provide 6 kJ? This is solved by simple division. Keeping in mind that 1 kJ = 1000 J, we haveNumber of photons * Energy per photon = 6 kJNumber of photons * (1.369*10^-17 J/photon) = 6 kJNumber of photons * (1.369*10^-17 J/photon) = 6000 JNumber of photons = 6000 J / (1.369*10^-17 J/photon)Number of photons = 4.382*10^20 photonsThis is how many photons (at this frequency) are needed to provide 6 kJ. How many moles of photons is this?Number of photons / number of photons in a mole = number of molesRecall that a mole of something is defined as 6.02*10^23of it. The same way a dozen eggs is defined as 12 eggs, a mole of eggs is 6.02*10^23 eggs. Equivalently, a mole of photons is 6.02*10^23 photons. SoNumber of photons / (6.02*10^23 photons per mole) = number of moles(4.382*10^20 photons) / (6.02*10^23 photons per mole) = number of moles7.279*10^-4 moles = number of molesForgive me if my arithmetic is off, as I don't have a good calculator handy. However, I believe this is the correct method to use.
How do you calculate how much energy in kJ do 3.0 moles of photons all with a wavelength of 655 nm contain?
The energy is 18,263.10e4 joules.
How do you calculate energy per mole of photons if you know joules per photon?
To calculate the energy per mole of photons from the energy per photon, you need to multiply the energy per photon by Avogadro's number (6.022 x 10^23) to account for the number of photons in a mole. The formula is: Energy per mole of photons = Energy per photon x Avogadro's number.
How much energy in kJ do 3.0 moles of photons all with a wavelength of 655 nm contain?
To calculate the energy of photons, you can use the equation E = hc/λ, where h is Planck's constant (6.626 x 10^-34 J·s), c is the speed of light (3.00 x 10^8 m/s), and λ is the wavelength. First, convert the wavelength to meters (655 nm = 655 x 10^-9 m). Plug the values into the equation to find the energy per photon, and then multiply by Avogadro's number to get the total energy for 3.0 moles of photons.
How much energy (in kJ) do 3.0 moles of photons all with a wavelength of 675 nm contain?
To calculate the energy of photons, we can use the formula (E = \frac{hc}{\lambda}), where (h) is Planck's constant ((6.626 \times 10^{-34} , \text{J s})), (c) is the speed of light ((3.00 \times 10^8 , \text{m/s})), and (\lambda) is the wavelength in meters (675 nm = (675 \times 10^{-9} , \text{m})). First, calculate the energy of one photon, then multiply by the number of moles (using Avogadro's number, (6.022 \times 10^{23} , \text{photons/mole})). Calculating this gives: Energy of one photon: [ E = \frac{(6.626 \times 10^{-34} , \text{J s})(3.00 \times 10^8 , \text{m/s})}{675 \times 10^{-9} , \text{m}} \approx 2.94 \times 10^{-19} , \text{J} ] Total energy for 3.0 moles of photons: [ \text{Total energy} = 3.0 , \text{moles} \times (6.022 \times 10^{23} , \text{photons/mole}) \times (2.94 \times 10^{-19} , \text{J}) \approx 5.34 \times 10^{5} , \text{J} ] Convert to kJ: [ 5.34 \times 10^{5} , \text{J} \div 1000 \approx 534 , \text{kJ} ] Thus, 3.0 moles of photons at 675 nm contain approximately 534 kJ of energy.
How many moles of photons with a wavelength of 660 nm are needed to provide 6 kJ of energy?
The energy carried by a photon is given byE = hfWhere h is Planck's constant (6.626x10^-34 Joule-seconds) and f is the frequency of the photon in Hertz (Hz).We are given the wavelength of the photon in the question in nanometers (nm). First, we need to convert this to (SI) units, because our equations only work with SI units. Then, we will calculate the frequency of the photon from its wavelength. Once we know the frequency of the photon we're interested in, we simply use the equation above to find the energy carried by one of them. Then we divide 6 kJ by that amount of energy, and the quotient will be the number of photons needed to carry 6 kJ. Finally, when we know the number of photons we need, we divide by the number of photons in a mole to get the number of moles.The SI unit of length is the meter (m). 1nanometer (nm) is 10^-9 meters.660 nm = 660 *10^-9 m = 6.6*10^-7 m.Now we will calculate this photon's frequency from its wavelength. These are related by the equationc = fLwhere c is the speed of light (3*10^8 m/s), f is the frequency of the photon and L is the wavelength of the photon.c = fL(3*10^8 m/s) = f * (6.6*10^-7 m)solving for f, we havef = (3*10^8 m/s) / (6.6*10^-7 m) = 4.54*10^15 s^-1Note that the unit of seconds (s) raised to the -1power is defined as 1 Hertz (Hz).f = 4.54*10^15 HzNow we will use the top equation to solve for the energy carried by one photon having this frequency.E = hfE = (6.626*10^-34 Js) * (4.54*10^15 Hz)E = 1.369*10^-17 JThis is how much energy is carried by one photon of wavelength 660 nm (which will also have a frequency of 4.54*10^15 Hz).How many of these do we need to provide 6 kJ? This is solved by simple division. Keeping in mind that 1 kJ = 1000 J, we haveNumber of photons * Energy per photon = 6 kJNumber of photons * (1.369*10^-17 J/photon) = 6 kJNumber of photons * (1.369*10^-17 J/photon) = 6000 JNumber of photons = 6000 J / (1.369*10^-17 J/photon)Number of photons = 4.382*10^20 photonsThis is how many photons (at this frequency) are needed to provide 6 kJ. How many moles of photons is this?Number of photons / number of photons in a mole = number of molesRecall that a mole of something is defined as 6.02*10^23of it. The same way a dozen eggs is defined as 12 eggs, a mole of eggs is 6.02*10^23 eggs. Equivalently, a mole of photons is 6.02*10^23 photons. SoNumber of photons / (6.02*10^23 photons per mole) = number of moles(4.382*10^20 photons) / (6.02*10^23 photons per mole) = number of moles7.279*10^-4 moles = number of molesForgive me if my arithmetic is off, as I don't have a good calculator handy. However, I believe this is the correct method to use.
How is stoichiometry used to calculate energy released when I mass of liquid freezes?
Stoichiometry can be used to calculate the energy released during the freezing of a liquid by calculating the moles of the liquid that freeze and then using the enthalpy of fusion of the substance (given in kJ/mol) to determine the total energy released during the process. The energy released can be found by multiplying the moles of liquid that freeze by the enthalpy of fusion value.
How to calculate excess moles of acid in titration?
To calculate the excess moles of acid in a titration, subtract the moles of base used from the initial moles of acid. This will give you the amount of acid that was not neutralized by the base and therefore the excess moles of acid present in the solution.
How do you calculate moles from molarity?
To calculate moles from molarity, you use the formula: moles = molarity x volume (in liters). Simply multiply the molarity of the solution by the volume of the solution in liters to find the number of moles present in the solution.
What would you need to do to calculate the molality of 0.2kg of nacl in 3kg of water?
To calculate the molality of a solution, you need to know the moles of solute and the mass of the solvent in kilograms. First, calculate the moles of NaCl in 0.2 kg: moles = mass (g) / molar mass. Then, calculate the molality by dividing the moles of solute by the mass of solvent in kg: molality = moles of solute / mass of solvent in kg.
calculate the number of moles in H2 molecules in 4g H2?
2 moles.
A solution contains 72.0 g of HCl and 468 g of (C6H6). What is the mole fraction of benzene?
First, calculate the moles of each component: moles of HCl = 72.0 g / molar mass of HCl and moles of C6H6 = 468 g / molar mass of C6H6. Then, calculate the total moles in the solution by adding the moles of each component. Finally, calculate the mole fraction of benzene by dividing the moles of C6H6 by the total moles in the solution.
How do you calculate moles of ions in 0.200 moles aluminum perchlorate.?
0,6 moles of (ClO4)3- and 0,2 mol Al
